LIBRARY 

OF  THE 

UNIVERSITY  OF  CALIFORNIA. 


GIFT    OF 


PROF,  W  B 

Class 


AN  INTRODUCTION  TO  THE  STUDY 


OP 


QUALITATIVE 


CHEMICAL  ANALYSIS. 


BY 


J.T.  McGILL,  Ph.D., 

Adjunct  Professor  of  Chemistry  in  Vanderbilt  University. 


PRINTED  FOR  THE  AUTHOR. 

PUBLISHING  HOUSE  OF  THE  M.  E.  CHUKCH,  SOUTH. 
J.  D.  BARBEE,  AGENT,  XASHVILLE,  TENN. 

1889. 


COPYKIOHT,  1889, 

BY 

J.  T.  McGiLL. 


PREFACE. 


THIS  little  book  is  intended  to  be  used  only  with  the  assistance 
of  a  teacher  and  as  an  introduction  to  some  full  work  on  qualita- 
tive analysis,  such  as  that  of  Fresenius  or  Prescott. 

Part  I.  contains  such  experiments  as  are  found  in  late  ele- 
mentary works  on  inorganic  chemistry.  These  experiments  are 
designed  to  train  the  student  in  the  construction  and  handling  of 
chemical  apparatus,  and  to  teach  him  how  to  observe  scientific- 
ally, reason  upon  his  observations,  and  draw  correct  conclusions 
from  them.  They  are  selected  for  the  most  part  with  special 
reference  to  the  analytical  work  that  follows  in  Part  II.,  and  in- 
clude the  preparation  of  many  of  the  reagents  used  in  qualita- 
tive analysis  and  the  examination  of  their  properties.  Part  I. 
may  be  omitted  by  students  who  have  taken  a  course  in  ele- 
mentary chemistry,  including  experiments  in  the  laboratory, 
such  as  Eliot  and  Storer's,  Remsen's,  or  Williams's. 

Part  II.  contains  the  special  feature  of  this  book — L  e.,  the  ap- 
plication to  qualitative  analysis  of  that  method  of  teaching  which, 
instead  of  imparting  directly  to  the  student  the  facts  or  laws 
which  are  the  object  of  the  study,  leads  him  by  questions  and 
suggestions  to  work  them  out  for  himself.  The  use  of  this 
method  for  two  years  in  Vanderbilt  University  has  led  me  to 
conclude  that  it  prevents  to  a  great  extent  the  blind  following  of 
directions,  which  the  student  is  so  liable  to  fall  into  in  qualita- 
tive analysis,  and  that  it  also  imparts  greater  interest  to  this 
study  by  requiring  the  student  to  test  by  practical  examples  the 
correctness  of  the  conclusions  which  he  has  arrived  at  as  the  re- 
sult of  his  own  investigations.  It  is  not  expected  that  the  stu- 
dent will  be  able  thus  to  construct  full  analytical  tables  which 

237406 


:    .»i  ^PREFACE. 


may  be  used  for  very  delicate  or  extensive  analysis ;  but  he  will 
become  prepared  to  use  such  tables  in  works  on  qualitative  anal- 
ysis, not  as  a  machine,  but  with  an  intelligent  understanding  of 
the  principles  involved. 

In  the  selection  and  description  of  experiments  I  have  used 
freely  the  works  of  Eliot  and  Storer,  Remsen,  Williams,  Hart, 
and  others. 

The  author  would  be  glad  to  have  the  benefit  of  corrections  or 
suggestions  that  may  occur  to  any  one  who  examines  or  uses  the 
book. 

Vanderbilt  University,  July,  1889. 


SUGGESTIONS  TO  TEACHERS. 

1.  Have  all  material  necessary  for  the  exercise  provided  be- 
forehand convenient  in  the  laboratory. 

2.  Discuss  in  the  class-room  with  the  students  the  work  of  the 
last  exercise ;  state  the  correct  results  and  inferences,  or  assign 
the  experiments  when  unsatisfactory  for  repetition,  giving  fuller 
instructions ;  give  directions  in  regard  to  the  experiments  of  the 
next  exercise ;  make  drawings  on  the  blackboard  of  apparatus 
to  be  used,  or  better  have  a  model  of  it  at  hand,  and  show  how 
it  is  constructed ;  indicate  the  method  of  performing  the  experi- 
ments, but  rarely  perform  the  experiments  yourself. 

3.  Caution  the  student  in  the  use  of  dangerous  chemicals  and 
the  performance  of  experiments  in  which  accidents  are  liable  to 
occur.    In  my  experience  more  accidents  have  happened  from 
the  careless  use  of  concentrated  sulphuric  acid  than  in  any  other 
way. 

4.  Require  the  pupils  at  the  close  of  each  exercise  to  deposit 
their  note-books  at  some  place  where  you  can  get  them  for  ex- 
amination.   A  case  of  "  pigeon-holes  "  numbered  to  correspond 
with  the  numbers  of  the  desks  will  be  found  convenient  as  a 
place  of  deposit  for  note-books,  and,  later  hi  the  course,  for  notes 
on  analytical  work  and  compounds  for  analysis. 


DIRECTIONS  TO  PUPILS. 


1.  Before  beginning  an  experiment  read  the  directions  through, 
and  provide  yourself  with  every  thing  which  you  can  foresee  to 
be  necessary  to  complete  the  experiment. 


EXPLANATIONS. 


2.  Direct  special  attention  to  constructing  neat  and  well-fitting 
apparatus,  and  keep  every  thing  about  your  desk  clean. 

3.  At  the  close  of  an  experiment  record  in  your  note-book  im- 
mediately (1)  the  material  employed,  a  brief  description  of  the 
apparatus  used,  the  method  of  procedure ;  (2)  what  takes  place  ; 
(3)  the  inferences  you  draw  or  what  you  have  learned  from  the 
experiment. 

4.  Leave  ample  space  in  your  note-book  for  making  corrections 
and  additions  without  erasing  any  of  the  first  record. 


EXPLANATIONS. 


1.  Dilute  acid  is  intended  unless  cone,  is  prefixed. 

2.  When  a  formula  is  given  without  the  name  of  the  com- 
pound the  name  and  formula  will  be  found  in  some  previous 
experiment. 

3.  An  interrogation  point  inclosed  in  parentheses  at  the  close 
of  a  statement  means  that  the  student  is  expected  to  determine 
by  experiment  whether  the  statement  be  true. 

4.  Weighed  quantities  of  substances  are  used  approximately, 
except  when  the  expression  "  weigh  out "  occurs. 


CONTENTS. 


PART  I.   EXPERIMENTS  IN  GENERAL  CHEMISTRY. 

EXPERIMENT  PAGE 

1.  Length 9 

2.  Volume 9 

3.  Weight 9 

4.  The  Bunsen  Burner , 10 

5.  Temperature  of  Ignition 10 

6.  The  Blow-pipe  and  Blast-lamp 11 

7.  Charcoal  as  a  Reducing  Agent 11 

8.  Sodium, Nitrate  as  an  Oxidizing  Agent 11 

9.  Heating  on  Asbestos  ;  on  Charcoal 12 

10.  Fusion  on  Platinum  Wire ;  on  Platinum  Foil 12 

11.  Heating  and  Bending  Glass 13 

12.  Evaporation,  Distillation,  Solution 13 

13.  Effect  of  Solutions  upon  Litmus 14 

14.  Dilute  Acids 15 

15.  Ammonia 15 

16.  Hydrochloric  Acid 16 

17.  Ammonium  Chloride 16 

18.  Salts  from  an  Acid  and  an  Hydroxide 17 

19.  Decomposition  of  Mercuric  Oxide 17 

20.  Oxides  and  Hydroxides  of  the  Alkali  Metals 18 

21.  Oxides  and  Hydroxides  of  the  Alkaline  Earth  Metals 18 

22.  Oxides  of  Magnesium,  Zinc,  Tin,  and  Aluminium 19 

23.  Salts  from  an  Acid  and  a  Metallic  Oxide 19 

24.  Sulphur 20 

25.  Oxides  of  Sulphur 20 

26.  Phosphorus  Pentoxide 20 

27.  Sodium  Sulphide 21 

28.  Iron  Sulphide ;  Hydrogen  Sulphide 22 

29.  Hydrogen  by  the  Action  of  a  Metal  on  Water 22 

30.  Hydrogen  by  the  Action  of  a  Metal  on  an  Acid 23 

31.  Oxygen 24 

32.  Nitrogen 25 


8  CONTENTS. 


EXPERIMENT  PAGE 

33.  Nitrous  Oxide 26 

34.  Nitric  Oxide 27 

35.  Nitric  Acid 27 

36.  Hydrochloric  Acid 28 

37.  Aqua  Eegia 29 

38.  Carbon  Dioxide. . .  .29 


PART  II.  EXPERIMENTS  IN  QUALITATIVE  ANALYSIS. 

THE  METALS. 

39.  Separation  of  the  Metals  into  Groups 31 

40.  Group  I.  Hydrochloric  Acid  Group 33 

41.  Group  II.  Hydrogen  Sulphide  Group 34 

42.  Group  III.  Ammonium  Sulphide  Group 37 

43.  Precautions  in  the  Separation  of  Groups  I.,  II.,  and  III ...  40 

44.  Group  IV.  Ammonium  Carbonate  Group 40 

45.  Group  V.  Sodium  Phosphate  Group * 42 

46.  Group  VI.  Soluble  Group 42 

47.  Separation  of  Groups  IV.,  V.,  and  VI 43 

48.  Some  Special  Cases  under  Group  III 44 

THE  ACIDS. 
Inorganic  Acids. 

49.  Barium  Chloride  Group 44 

50.  Silver  Nitrate  Group 45 

51.  Silicic  Acid 45 

52.  Hydrofluoric  Acid 46 

53.  Boric  Acid 46 

54.  Sulphurous  and  Hyposulphurous  Acids 46 

55.  Chloric  Acid 47 

Organic  Acids. 

56.  Test  of  an  Organic  Compound 47 

57.  Cone.  Sulphuric  Acid  on  Organic  Acids 47 

58.  Oxalic,  Tartaric,  Citric,  and  Malic  Acids 48 

59.  Succinic,  Benzoic,  Formic,  and  Acetic  Acids 48 

60.  Organic  Compounds  Insoluble  in  Water 48 


PART  I. 

EXPERIMENTS  IN  GENERAL  CHEMISTRY. 


LENGTH. 

1.  Estimate  by  the  eye  1  dm.     Close 
the  book  and  lay  off  this  distance  on  a 
slip  of  paper;  divide  it  into  cm.;  divide 
1  cm.  into  mm.     Compare  the  scale  you 
have  made  with  the  model.     Repeat  this 
experiment  several  times.     Estimate  the 
dimensions   of  glass   tubing,  test-tubes, 
etc. 

VOLUME. 

2.  Pour  25  cc.  of  water  into  a  small 
beaker.     Mark  the  level  of  the  water  by 
means  of  a  gummed  label.     Graduate  a 
larger  beaker  to  100,  200,  and  500  cc., 
and  a  test-tube  to  5,  10,  and  25  cc.     Es- 
timate the  capacity  of  different  vessels 
at  your  desk. 

WEIGHT. 

3.  (a)  Weigh    the    small    graduated 
beaker.     Pour  into  it  25  cc.  of  distilled 
water  and  weigh.     Pour  out  the  water, 
wipe    the    beaker    perfectly    dry,    and 
weigh   in  it   25   cc.  of  cone,  sulphuric 
acid.     What  is  the  specific   gravity  of 
cone,   sulphuric   acid?     How  many  cc. 


10  EXPERIMENTS   IN 


cone,  sulphuric  acid  weigh  25  g.  ?     How  many  cc.  of  al- 
cohol, specific  gravity  0.925,  weigh  50  g.  ? 

(6)  Balance  two  filter  papers  on  the  scales.  Weigh 
out  1  g.  of  sodium  chloride  (common  salt),  1  g.  of  lead 
chloride  (PbCl2),  1  g.  of  calcium  phosphate  (Ca3PO4). 
By  subdividing  these,  estimate  the  amount  of  each  that 
would  weigh  0.5  g.,  0.1  g.  Preserve  the  1  g.  weights 
of  the  above  substances  for  future  use. 

THE  BUNSEN  BURNER. 

4.  (a)  Light  and  Heat  of  Flame. — Examine  the  struct- 
ure of  a   Bunsen  burner.     Compare  the  flame  of  the 
burner  when  the  valve  is  open  with  the  flame  when  the 
valve  is  closed.     In  which  flame  will  a  platinum  wire 
glow  more  brightly?     Which  flame  gives  more  light? 
which  more  heat?    A  deposit  of  carbon  is  made  on  a 
piece  of  porcelain  held  in  one  of  these  flames  (  ?  ).     Can 
you  burn  the  carbon  off?     Hub  together  two  pieces  of 
charcoal  so  that  a  fine  dust  falls  into  the  colorless  flame. 
From  the  above  experiments  what  inferences  could  you 
draw  in  regard  to  the  causes  of  light  and  heat  in  flame  ? 

(ft)  Form  of  Flame. — Make  drawings  of  the  flames  of 
the  Bunsen  burner.  Hold  a  platinum  wire  across  the 
flame  (try  both  flames)  at  different  heights  above  the 
burner.  Try  a  similar  experiment  with  a  splinter ;  with 
a  glass  rod.  Thrust  the  sulphur  end  of  a  match  quickly 
into  the  center  of  the  flame.  Hold  a  piece  of  paper  hor- 
izontally in  the  flame  and  withdraw  it  just  before  it 
would  ignite.  Use  wire  gauze  instead  of  paper,  omit- 
ting the  withdrawal.  What  do  you  learn  from  these  ex- 
periments ? 

TEMPERATURE  OF  IGNITION. 

5.  Pour  1  cc.  of  carbon  bisulphide  (CS2)  into  a  porcelain 
crucible.     Can  you  ignite  it  by  thrusting  a  heated  glass  or 


GENERAL  CHEMISTRY.  11 

iron  rod  into  the  mouth  of  the  crucible  ?  Bring  the  heated 
rod  in  contact  with  a  little  sulphur.  Can  you  ignite  coal 
gas  with  the  heated  rod  ?  Suspend  a  spiral  of  platinum 
wire  in  a  small  flame  above  a  burner.  When  the  plat- 
inum is  at  a  white  heat  put  out  the  flame  with  a  sudden 
puff  of  air.  Explain  the  result.  Lower  a  piece  of  wire 
gauze  horizontally  into  the  flame  of  a  burner.  Can  you 
thus  extinguish  a  small  flame  ?  Hold  a  piece  of  wire 
gauze  horizontally  three  or  four  cm.  above  a  burner; 
turn  on  the  gas  and  try  to  light  it  above  the  gauze.  If 
you  succeed,  gradually  raise  the  gauze.  What  is  the 
office  of  a  wire  gauze  placed  under  vessels  that  are  being 
heated  ? 

THE  BLOW-PIPE  AND  BLAST-LAMP. 

6.  Examine  the  structure  of  the  blow-pipe ;  the  blast- 
lamp.     Compare   their  flames  with   the   burner  flame. 
Practice  with  the  blow-pipe  until  you  can  blow  a  steady 
stream  of  air  for  two  minutes. 

CHARCOAL  AS  A  REDUCING  AGENT, 

7.  Mix  in  a  mortar  1  g.  of  lead  carbonate  with  a  suf- 
ficient quantity  of  powdered  charcoal  to  nearly  fill  a 
small  porcelain  crucible.     Cover  the  crucible  and  heat 
it  over  the  blast-lamp  for  three  minutes.     Pour  the  con- 
tents of  the  crucible  into  a  mortar,  rub  with  water,  pour 
off  the  charcoal  and  water  and  look  for  scales  of  lead  in 
the  bottom  of  the  mortar. 

SODIUM  NITRATE  (NaN03)  AS  AN  OXIDIZING  AGENT. 

8.  Fill  a  porcelain  crucible  half-full  of  XaNO3  and  heat 
until  the  salt  fuses.     Throw  powdered  charcoal  in  small 
quantities  upon  the  fused  mass.     The  charcoal  is  oxid- 


12  EXPERIMENTS   IN 


ized,  and  burns,  forming  carbon  dioxide  gas.     The  chem- 
ical action  that  takes  place  may  be  expressed  as  follows : 

Carbon  (charcoal)  and  oxygen  form  carbon  dioxide. 

C      +      02      =      CO,. 

Potassium    nitrate    (JOTO3)    and    potassium    chlorate 
(KC1O3)  are  likewise  used  as  oxidizing  agents. 

HEATING  ON  ASBESTOS;  ON  CHARCOAL. 

9.  (a)  Place  0.1  g.  of  lead  carbonate  on  a  slip  of  asbes- 
tos, and  heat  it  in  the  yellow  or  in  the  inner  flame  of 
the  blow-pipe  until  a  globule  of  lead  is  obtained.     Heat 
the  globule  in  the  point  of  the  blow-pipe  flame.     Oxide 
of  lead  is  formed.     Eeduce  the  oxide  to  lead.     Similar 
experiments  may  be  tried  with  Zn,  Sn,  Cu,  Sb,  Bi,  or 
their  compounds. 

(£>)  Try  to  reduce  some  of  these  compounds  to  the  me- 
tallic state  by  heating  them  on  a  piece  of  charcoal  be- 
fore the  blow-pipe.  Is  the  reduction  effected  more  easily 
on  asbestos  or  charcoal  ? 

FUSION  ON  PLATINUM  WIRE;  ON  PLATINUM  FOIL. 

Borax,  sodium  metaphosphate,  and  sodium  carbonate 
are  often  used  as  fluxes  in  heating  substances  on  plat- 
inum wire,  platinum  foil,  charcoal,  or  in  crucibles.  Ni- 
trates of  sodium  or  of  potassium  are  sometimes  used 
mixed  with  these.  Platinum  must  not  be  used  when  the 
substance  may  be  reduced  to  a  metallic  state. 

10.  (a)  Make   a   transparent   sodium   metaphosphate 
bead  in  a  loop  on  the  end  of  a  platinum  wire,  bring  it 
while  hot  in  contact  with  a  small  quantity  of  powdered 
manganese  carbonate,  heat  in  the  oxidizing,  then  in  the 
reducing  flame  of  either  the  Bunsen  burner  or  the  blow- 
pipe.    Observe  the  change  in  color.     Try  a  similar  ex- 
periment with  a  borax  bead  and  cobalt  carbonate. 


GENERAL  CHEMISTRY.  13 

(6)  Mix  0.05  g.  of  manganese  carbonate  with  0.2  g.  of 
sodium  carbonate  and  heat  on  a  platinum  foil. 

HEATING  AND  BENDING  GLASS. 

The    operations    should    first   be   performed   by  the 
teacher  in  the  presence  of  the  class. 

11.  Eound  the  ends  of  some  glass  rods  in  the  hottest 
part  of  the  flame  of  the  Bunsen  burner.     Make  half  a 
dozen  ignition  tubes  by  melting  in  two  at  the  middle 
pieces  of  glass  tubing  about  15  cm.  long.     Bend  eight 
tubes  at  right  angles,  making  six  of  them  with  the  limbs  6 
and  9  cm.,  one  25  and  10  cm.,  and  one  35  and  8  cm.  long. 
Round  the  sharp  edges  of  the  ends  of  these  in  the  flame. 

EVAPORATION,  DISTILLATION,  SOLUTION. 

12.  (a)  Dissolve  1  g.  NaCl  in  20  cc.  water.     Evaporate 
half  of  the  solution  to  dryness  in  a  porcelain  dish.    What 
remains  ?    Taste  it.    Distill  the  remainder  of  the  solution, 
heating  carefully  until  the  distillate  amounts  to  several 
cc.     (The  distillate  is  the  part  that  passes  over.)     Evap- 
orate the  distillate  to  dryness.     What  remains  ?    A  con- 
venient apparatus  for  this  purpose  may  be  made  by  con- 
necting two  large  test-tubes  by  a  tube  15  cm.  long,  curved 
so  that  the  test-tubes  set  at  about  right  angles  to  each 
other.     A  groove  in  the  side  of  the  stopper  of  the  tube 
used  to  collect  the  distillate  will  permit  the  escape  of  un- 
condensed  vapor. 

(6)  Place  0.1  g.  powdered  barium  sulphate  (BaSO4)  in  a 
test-tube.  Pour  into  the  test-tube  1  cc.  water.  Does 
the  BaSO^  appear  to  dissolve?  Add  5  cc.  water,  boil, 
and  filter.  Evaporate  a  portion  of  the  filtrate — z.e.,  the 
part  that  runs  through.  Is  BaSO4  soluble  in  water  ? 

(c)  Weigh  out  0.1  g.  calcium  sulphate  (CaSO4).  Shake 
it  up  with  5  cc.  water,  and  let  it  stand  a  few  minutes.  If 


14  EXPERIMENTS  IN 


it  does  not  all  dissolve,  add  more  water,  shake,  and  let 
it  stand.  Determine  in  how  many  times  its  weight  of 
cold  water  CaSO4  is  soluble,  1  cc.  water  weighing  1  g. 
How  could  the  exact  solubility  of  CaSO4  be  determined 
without  adding  a  sufficient  amount  of  water  to  dissolve 
the  whole  of  the  weighed  amount  taken  ?  Save  10  cc. 
of  the  solution  for  12  (e)  and  12  (/). 

(d)  Make  a  strong  solution  of  0.2  g.  powdered  crystal- 
line lead  chloride  in  hot  water.  Cool  the  solution.  Explain 
the  result.     Does  all  of  the  lead  chloride  crystallize  out  ? 

(e)  Add  10  cc.  alcohol  to  5  cc.  CaSO4  solution. 

(/)  Add  10  cc.  alcohol  to  a  strong  solution  of  sodium 
chloride,  filter,  and  dry  the  precipitate.  Is  it  sodium 
chloride  ? 

(<7)  Add  cone,  hydrochloric  acid  to  a  strong  solution 
of  sodium  chloride,  filter,  and  evaporate  off  the  acid 
from  the  residue.  What  is  this  residue  ? 

(h)  The  same  as  (#),  using  barium  chloride  instead  of 
sodium  chloride. 

(i)  Dissolve  1  g.  of  calcium  phosphate  (bone  ash)  in 
hydrochloric  acid,  and  evaporate  to  dryness  until  the  res- 
idue does  not  smell  of  acid.  Is  calcium  phosphate  at- 
tacked by  the  acid,  or  simply  held  in  solution  ? 

(j)  Add  solid  sodium  carbonate  to  1  cc.  hydrochloric 
acid  until  effervescence  ceases.  Is  the  gas  that  comes 
off  combustible?  Evaporate  the  solution  to  dryness. 
Compare  the  taste  of  the  residue  with  that  of  the  sodi- 
um carbonate.  How  does  this  solution  differ  from  the 
preceding  ? 

(A:)  Dissolve  sodium  carbonate  in  water,  evaporate, 
and  examine  the  residue. 

EFFECT  OF  SOLUTIONS  UPON  LITMUS. 

13.  Test  the  reaction  of  all  the  solutions  at  your  desk 


GENERAL  CHEMISTRY.  15 

to  red  and  to  blue  litmus  paper.  Those  turning  blue  to 
red  are  said  to  have  an  acid  reaction,  those  turning  red 
to  blue  an  alkaline  reaction,  and  those  not  changing  the 
color  of  either  red  or  blue  litmus  a  neutral  reaction. 

DILUTE  ACIDS. 

14.  (a)  Pour  5  cc.  cone,  sulphuric  acid  into  25  cc.  wa- 
ter.   Notice  the  rise  in  temperature.    This  is  the  strength 
of  the  dilute  sulphuric  acid  at  your  desk.     Water  must 
never  be  poured  into  cone,  sulphuric  acid.     Why  not  ? 

(6)  Pour  5  cc.  cone,  hydrochloric  acid  into  25  cc.  wa- 
ter. This  is  the  strength  of  the  dilute  hydrochloric  acid 
at  your  desk. 

(c)  Pour  5  cc.  cone,  nitric  acid  into  25  cc.  water.  This 
is  the  strength  of  the  dilute  nitric  acid  at  your  desk. 

(<2)  Dilute  with  water  5  cc.  of  each  of  the  above  solu- 
tions (a),  (6),  (c),  to  500  cc.,  and  compare  the  acidity  by 
tasting. 

AMMONIA. 

15.  (a)  Distill  some  sodium   hydrate  solution.     Test 
the  distillate  with  litmus  paper ;  the  residue. 

(6)  Pour  25  cc.  cone,  ammonia  into  a  small  flask, 
through  the  stopper  of  which  is  inserted  a  straight  glass 
tube  25  cm.  long,  open  at  both  ends.  Hang  a  test-tube  or 
small  flask  over  the  end  of  the  tube.  Warm  the  ammo- 
nia. When  the  test-tube  is  full  of  the  gas,  which  may 
be  known  by  the  smell  of  ammonia  at  its  mouth, 
take  it  off,  keeping  its  mouth  downward,  and  dip  the 
end  of  it  under  the  surface  of  the  water.  Explain  what 
happens.  Is  ammonia  gas  combustible  ?  Is  it  lighter  or 
heavier  than  air  ?  What  is  the  color  of  the  gas  ?  Is  it 
more  soluble  in  cold  or  hot  water  ?  Save  100  cc.  of  the 
gas  in  a  closely-stoppered,  inverted  bottle  for  17  (a).  A 
gas  is  said  to  be  collected  by  upward  displacement  when 


16  EXPERIMENTS  IN 


collected  as  in  this  experiment.  Ammonia  gas  has  the 
formula  NH3.  Its  solution  in  water  is  called  ammonia 
(also  ammonium  hydroxide  and  ammonium  hydrate),  and 
has  the  formula  NH4OH ;  but  in  reactions  it  is  often 
written  NH8. 

HYDROCHLORIC  ACID. 

16.  Heat  25  cc.  of  cone,  hydrochloric  acid  in  a  small 
flask  under  a  hood.     A  gas  is  given  off.     Collect  it  by 
letting  the  outlet  tube  extend  downward  almost  to  the 
bottom  of  the  receiving  vessel  which  may  be  covered  by 
a  piece  of  paper.     A  piece  of  moistened  blue  litmus  pa- 
per at  the  point  where  the  outlet  tube  pierces  the  paper 
covering  will  indicate  when  the  vessel  is  filled  by  turning 
red.     The  gas  is  thus  collected  by  downward  displace- 
ment.    Fill  several  bottles  and  cover  their  mouths  with 
watch-glasses  or  glass  plates.     Invert  one  of  the  bottles 
in  water,  and  remove  the  glass  covering.     Is  the  gas 
more  soluble  in  cold  or  hot  water  ?     Is  it  combustible  ? 
Is  it  lighter  or  heavier  than  air?     Hydrochloric  acid  gas 
has  the  formula  HC1.     Its  solution  in  water  is  called  hy- 
drochloric acid,  or  muriatic  acid,  and  in  reactions  is  also 
written  HC1. 

AMMONIUM  CHLORIDE. 

17.  (a)  Bring  the  mouth  of  the  bottle  of  ammonia  gas 
saved  from  experiment  15  over  a  bottle  of  hydrochloric 
acid  gas,  and  mix  the  gases. 

(6)  Evaporate  to  dryness  10  cc.  of  ammonium  chloride 
solution,  and  compare  the  residue  with  the  solid  product 
obtained  by  mixing  hydrochloric  acid  gas  and  ammonia 
gas,  (1)  by  tasting,  (2)  by  heating  on  a  piece  of  platinum 
foil. 

(c)  Suspend  a  drop  of  concentrated  hydrochloric  acid 
on  the  end  of  a  glass  rod,  and  bring  it  near  the  mouth  of 
a  bottle  of  ammonia.  What  is  the  smoke  that  is  formed  ? 


GENERAL  CHEMISTRY.  17 

Ammonia  gas  and  hydrochloric  acid  gas  form  ammonium  chloride. 

HC1      = 


Bring  a  drop  of  the  acid  near  a  bottle  of  sodium  hydrox- 
ide. (The  hydroxides  are  also  called  hydrates.)  Try 
the  effect  of  ammonia  gas  on  other  acids  in  the  same 
way,  or  by  bringing  a  drop  of  cone,  ammonia  near  the 
mouth  of  the  bottle  of  acid. 

SALTS  FROM  AN  ACID  AND  AN  HYDROXIDE. 

18.  (a)  Mix  till  neutral  to  litmus  paper  dilute  HC1 
with  10  cc.  dilute  NH4OH.     If  too  much  acid  be  added, 
neutralize  the  acid  by  adding  drop  by  drop  very  dilute 
NH4OH.     Evaporate  the  neutral  solution  to  dryness; 
taste  the  residue  ;   heat  it  to  redness.     The  substance 
will  be  recognized  as  NH4C1.     The  reaction  may  be  writ- 

ten thus  :  Water. 

NH4OH  +  HC1  =  NH4C1  +  H20. 

(6)  Try  some  of  the  following  experiments,  and  write 
the  reaction  for  all  of  them  :  HC1  neutralized  with  potas- 
sium hydroxide  (KOH),  with  sodium  hydroxide  (NaOH)  ; 
nitric  acid  (HNO3),  sulphuric  acid  (H2SO4),  phosphoric 
acid  (H3PO4),  each  neutralized  with  KE4OH,  KOH, 
NaOH,  respectively.  By  neutralizing  an  acid  with  an  hy- 
droxide a  salt  is  formed. 

DECOMPOSITION  OF  MERCURIC  OXIDE. 

19.  Heat  1  g.  of  mercuric  oxide  (HgO)  in  a  test-tube. 
Insert  a  glowing  splinter  into  the  mouth  of  the  tube. 
The  gas  given  off,  which  causes  the  glowing  splinter  to 
ignite,  is  oxygen.    Mercuric  oxide  is  decomposed  by  heat- 
ing into  mercury  and  oxygen.     HgO  =  Hg  -j-  O.     Ox- 
ygen was  first  obtained  in  this  way.     Some  other  oxides 
are  decomposed  by  heating  into  the  metal  and  oxygen. 


18  EXPERIMENTS   IN 


OXIDES  AND  HYDROXIDES  OF  THE  ALKALI  METALS. 

20.  (a)  Place   in  a  small  porcelain   dish  a   piece  of 
potassium  the  size  of  a  pea.     The  change  of  color  at  the 
surface  on  remaining  in  the  air  is  due  to  the  formation 
of  the  oxide  of  potassium.     (Are  gold,  silver,  tin,  zinc, 
iron  oxidized  thus  ?)     Heat  gently.     The  oxidation  takes 
place  more  rapidly.     After  the  potassium  is  completely 
oxidized  allow  it  to  cool,  and  then  add  a  few  drops  of 
water.     What  causes  the  rise  in  temperature  ?     Test  the 
solution  with  litmus  paper.     What  doesihe  solution  con- 
tain ?     K2O  +  H2O  =  ? 

(6)  Perform  a  similar  experiment  with  sodium. 
Na2O  -f  H2O  =  ?  The  teacher  should  furnish  for  ex- 
periments in  which  they  are  required  K,  Na,  and  P  al- 
ready cut  into  pieces  of  the  proper  size. 

OXIDES  AND   HYDROXIDES  OF  THE  ALKALINE 
EARTH    METALS. 

21.  (a)  Is  calcium  carbonate  (CaCO3)  soluble  in  water? 
Try  a  particle  of  it.     Test  the  liquid  with  litmus  paper. 
Heat  5  g.  CaCO3  in  an  open  porcelain  crucible  over  a 
blast-lamp  for  several  minutes.     CaCO3  =  CaO  -f  CO2. 

Pour  a  portion  of  the  contents  of  the  crucible  into 
a  porcelain  dish,  and  when  that  in  the  dish  has  cooled 
pour  on  it  a  little  water.  What  causes  the  heat  ?  Cal- 
cium hydroxide  is  formed:  CaO  +  H2O  =  Ca(OH)2. 
Test  with  litmus  paper.  The  solution  of  Ca(OH),  in 
water  is  called  lime-water.  Prove  that  Ca(OH)2  is  sol- 
uble in  water.  Is  it  more  soluble  than  KOH  or  NaOH  ? 

(7?)  Hub  some  moist  Ca(OH)2  with  a  particle  of  any 
ammonium  salt.  What  is  given  off?  Write  the  reac- 
tion. Heat  moist  Ca(OH)2  with  an  ammonium  salt. 
Ammonia  gas  for  making  ammonium  hydroxide  (ammo- 
nia) is  thus  manufactured.  Will  KOH  or  NaOH  de- 


GENERAL  CHEMISTRY.  19 

compose  an  ammonium  salt  ?     How  can  you  distinguish 
an  ammonium  compound  from  other  compounds  ? 

OXIDES  OF   MAGNESIUM,  ZINC,  TIN,  AND 
ALUMINIUM. 

22.  Heat    a    platinum   wire.     There   is   no    chemical 
change.     Heat  a  magnesium  wire.     Tilt  the  burner  to 
one  side  so  that  the  magnesium  oxide  formed  in  the  com- 
bustion may  be  caught  in  a  dish.     Moisten  the  oxide 
with  water  and  test  with  litmus  paper.     Is  it  soluble  in 
water?  in  hydrochloric  acid?     Does  hydrochloric  acid 
form  a  new  compound  with  it  or  merely  hold  it  in  solu- 
tion as  in  the  case  of  calcium  phosphate  in  12  (/)  ?     Is 
iron  easily  oxidized  ?     Does  oxide  of  iron  part  with  ox- 
ygen so  easily  as  HgO  ?     Is  it  soluble  in  water  ?     What 
would  you  say  of  ZnO,  SnO2,  A12O3  in  these  respects? 

It  is  seen  from  experiments  20  and  21  that  the  oxides 
of  the  alkali  metals  and  of  the  alkaline  earth  metals  unite 
with  water  to  form  soluble  hydroxides.  The  oxides  of 
most  of  the  other  metals  do  not  combine  directly  with  wa- 
ter, but  under  certain  conditions  they  form  hydroxides 
almost  insoluble  in  water.  The  hydroxides  of  the  metals 
are  called  bases.  They  unite  with  acids  to  form  salts  (18). 

SALTS  FROM  AN  ACID  AND  A  METALLIC  OXIDE. 

23.  (a)  Dissolve    1   g.  of  zinc  oxide  in   dilute   HC1. 
Evaporate  to  dryness.     The  residue  is  a  salt  soluble  in 
water.     ZnO  -f  2HC1  =  ZnCl2  +  H2O.     By  dissolving  a 
metallic  oxide  in  an  acid  a  salt  is  formed. 

(6)  Dissolve  copper  oxide  (CuO)  in  H2SO4.  Write  the 
reaction.  Evaporate  till  a  drop  of  the  solution  placed 
on  a  watch-glass  will  deposit  crystals  on  cooling.  Set 
the  solution  aside  to  cool. 

(c)  Heat  in  a  crucible  or  dish  1  g.  crystalline  copper 
sulphate.  CuSO4,  5H2O  =  CuSO4  +  5H2O. 


20  EXPERIMENTS  IN 


SULPHUR. 

24.  (a)  Heat  slowly  15  g.  S  in  a  test-tube.     Continue 
to  heat  until  it  boils,  noting  the  changes  which  the  S 
undergoes.     Pour  the  melted  S  in  a  fine  stream  into  a 
dish  of  water.     Examine  the  S  that  has  been  poured 
into  the  water.     Will  it  retain  its  plastic  condition  if 
kept  under  water  ?  in  the  open  air  ? 

(6)  Melt  in  a  crucible  or  in  a  small  beaker  25  g.  roll 
sulphur.  Let  it  cool  slowly  and  quietly.  When  a  thin 
crust  has  formed  on  the  surface,  make  a  hole  in  the  crust 
and  pour  oat  the  liquid  S.  Examine  the  crystals  that 
remain  in  the  vessel. 

(c)  Pour  3  cc.  carbon  bisulphide  (CS2)  upon  1  g.  S  in 
a  test-tube.     Close  the  mouth  of  the  test-tube  and  shake 
for  a  few  moments.     When  the  S  has  dissolved  pour  the 
solution  into  a  watch-glass.     Examine  the  crystals  that 
form  as  the  CS2  evaporates. 

(d)  Add  a  few  drops  of  cone.  HC1  to  5  cc.  ammonium 
sulphide.     Sulphur  separates  from  the  liquid  in  a  finely- 
divided  state. 

OXIDES  OF  SULPHUR. 

25.  Make  a  hollow  in  the  end  of  a  piece  of  chalk. 
Place  in  it  some  sulphur.     Ignite  the  sulphur  and  sus- 
pend it  in  a  large  bottle.     The  odor  is  that  of  sulphurous 
anhydride   formed   by  the   oxidation   of   the    sulphur. 
S  -f  O2  =SO2.     Pour  some  water  into  the  bottle.     Test 
the  solution  with  litmus  paper.     SO2  is  an  acid-forming 
oxide.     The  acid  formed  is  sulphurous  acid.     SO2  +  H2O 
=  H2SO3.     A  higher  oxide  of  sulphur  (SO3)  forms  with 
water  sulphuric  acid.     H2O  -f  SO3  =  H2SO4. 

PHOSPHORUS  PENTOXIDE. 

26.  Eepeat  experiment  25,  using  instead  of  S  a  piece 
of  dry  phosphorus  half  the  size  of  a  pea.     The  white  va- 


GENERAL  CHEMISTRY.  21 

por  is  phosphorus  pentoxide  (P2O5).  With  water  it  forms 
phosphoric  acid.  P2O5  +  3H2O  =  2H3PO4.  What  salt 
is  formed  by  neutralizing  this  acid  with  XaOH  ?  with 
Ca(OH)2  ?  Write  the  reactions. 

The  oxides  of  most  of  the  non-metallic  elements  are 
acid-forming  oxides — i.  e.,  unite  with  water  to  form  acids 
(25  and  26). 

SODIUM  SULPHIDE. 

Sulphur  forms  many  compounds  analogous  to  those  of 
oxygen.  Recurring  to  experiment  20,  perform  the  fol- 
lowing experiments  and  write  the  reactions : 

27.  (a)  Cover  a  piece  of  sodium  in  a  porcelain  crucible 
with  sulphur,  put  on  the  crucible  lid,  and  warm  gently. 
When  the  crucible  has  cooled  dissolve  a  portion  of  its 
contents  in  a  few  drops  of  water,  reserving  the  remainder 
for  27  (c).  Test  the  solution  with  litmus  paper.  The 
solution  contains  sodium  sulphide  (Na2S).  Place  a  drop 
of  it  on  a  silver  coin.  The  dark  substance  is  silver  sul- 
phide (Ag2S).  Eub  the  spot  with  a  little  potassium  cy- 
anide. (Potassium  cyanide  is  very  poisonous.)  Add  a 
few  drops  of  the  JS"a2S  solution  to  a  solution  of  silver  ni- 
trate in  a  test-tube,  then  gradually  add  potassium  cyanide 
solution. 

(6)  Try  the  effect  of  barium  sulphate  on  a  silver  coin ; 
of  sodium  carbonate.  Heat  on  charcoal  before  the  blow- 
pipe a  mixture  of  sodium  carbonate  and  barium  sulphate. 
Test  the  melted  mass  as  the  product  of  sodium  and  sul- 
phur was  tested  in  27  (a).  How  can  the  presence  of  sul- 
phur be  detected  in  an  insoluble  compound  ? 

(c)  Add  HC1  to  the  sodium  sulphide  reserved  from  27 
(a)  until  the  solution  is  neutral  to  litmus.  Compare  with 
18  and  write  the.  reaction.  The  gas  given  off  is  hydro- 
gen sulphide  (H2S).  What  remains  in  solution? 


22  EXPERIMENTS  IN 


IRON  SULPHIDE;    HYDROGEN  SULPHIDE. 

28.  (a)  Examine  fine  iron  filings  and  powdered  sulphur 
as  follows :  (1)  With  a  magnifying-glass,  (2)  by  bring- 
ing a  magnet  near  them,  (3)  solubility  in  water,  (4)  sol- 
ubility in  carbon  bisulphide,  (4)  solubility  in  acid. 

(ft)  Mix  thoroughly  2  g.  of  iron  filings  with  2  g.  of 
sulphur.  Examine  small  portions  of  the  mixture  as  in 
28  (a). 

(c)  Pour  the  mixture  into  a  dry  test-tube,  and  heat  till 
the  mixture  glows.     When  the  glowing  has  ceased  take 
out  the  mass,  and  examine  small  portions  as  in  28  (a). 

(d)  Put  the  remainder  of  the  iron  sulphide  into  a  test- 
tube,  add  H2SO4,  and  test  the  properties  of  the  gas. 
What  remains  in  solution  in  the  test-tube?     Write  the 
reaction.     The  solution  of  the  gas  in  water  is  the  hydro- 
gen sulphide  at  your  desk.     The  solution  of  ammonium 
sulphide  at  your  desk  is  made  by  saturating  a  solution 
of  ammonia  with  hydrogen  sulphide  gas,  or  a  solution  of 
hydrogen  sulphide  with  ammonia  gas. 

(e)  Try  the  effect  of  acids,  dilute  and  strong,  upon  hy- 
drogen sulphide  solution ;  upon  ammonium  sulphide  so- 
lution. 

(/)  Try  the  effect  of  a  few  drops  of  hydrogen  sul- 
phide on  a  solution  of  silver  nitrate ;  lead  acetate.  Try 
ammonium  sulphide  on  these  solutions. 

HYDROGEN  BYTHE  ACTION  OFAMETALON  WATER. 

29.  (a)  Throw  a  piece  of  potassium  the  size  of  a  grain 
of  wheat   on  water.     Notice   the   color   of  the   flame. 
What  burns  ?     Test  the  water  with  litmus  paper. 

(ft)  Throw  a  piece  of  sodium  as  large  as  a  pea  on  some 
water  in  a  porcelain  crucible.  Wrap  a  piece  of  sodium 
in  a  piece  of  thin  paper  and  throw  it  on  the  water  in  the 
crucible.  Test  the  water  with  litmus  paper.  Evaporate 


GENERAL  CHEMISTRY.  23 

to  dryness  and  compare  the  residue  with  that  obtained 
by  evaporating  some  solution  of  sodium  hydroxide. 
Neutralize  with  HC1.  What  is  formed?  Did  the  so- 
dium get  oxygen  from  the  air  in  burning  on  the  water  ? 

(c)  Fill  a  test-tube  with  water  and  invert  it  in  a  dish 
of  water.  Take  care  that  no  bubble  of  air  remains  in 
the  tube.  At  the  close  of  the  experiments  with  hydro- 
gen (30)  explain  the  reason  of  this  precaution.  Now 
take  a  small  piece  of  dry  sodium  on  the  point  of  a  rat- 
tail  file  and  thrust  it  under  the  mouth  of  the  test-tube  so 
that  the  sodium  will  rise  in  it.  Close  the  mouth  of  the 
test-tube  with  the  thumb,  and  open  it  near  a  flame.  The 
gas  that  burns  is  hydrogen.  Test  the  water  in  the  dish 
with  litmus  paper.  What  elements  does  this  experiment 
prove  water  to  contain  ? 

Water  is  decomposed  by  iron  at  a  high  temperature. 
Water  may  be  decomposed  by  a  current  of  electricity, 
and  the  volume  of  hydrogen  and  oxygen  measured.  The 
volume  of  the  hydrogen  is  then  found  to  be  twice  that 
of  the  oxygen.  The  action  of  an  acid  on  a  metallic  oxide 
has  been  found  to  result  in  the  formation  of  a  salt  and 
water  (25)— e.  g.,  ZnO  +  H2SO4  =  ZnSO4  +  H2O.  Let 
us  try  the  effects  of  an  acid  on  a  metal. 

HYDROGEN    BY  THE  ACTION    OF  A    METAL    ON   AN 
ACID. 

30.  Pour  H2SO4  on  some  pieces  of  Zn  in  a  small 
flask  or  large  test-tube.  Collect  the  gas  over  water  in 
test-tubes  at  first,  and  ignite  the  gas  as  in  29  (c).  What 
do  you  infer  from  the  way  the  gas  first  collected  burns  ? 
What  precautions  must  be  observed  in  making  this  gas  ? 
When  the  gas  has  been  coming  off  for  several  minutes, 
and  when  collected  in  a  test-tube  is  found  to  burn  qui- 
etly, collect  several  bottles  of  the  gas  and  try  the  follow 
ing  experiments : 


24  EXPERIMENTS  IN 


(a)  Thrust  a  lighted  taper  into  a  bottle  of  the  gas 
held  mouth  downward  (why  not  mouth  upward  ?) ;  with- 
draw the  taper ;  insert  again  while  the  gas  is  burning. 

(6)  Fill  a  small  bottle  with  air  and  hydrogen  in  the 
proportion  of  three  to  one  by  volume.  Let  the  mixture 
stand  for  awhile,  and  then  bring  the  mouth  of  the  bottle 
near  a  flame.  How  could  hydrogen  be  emptied  from 
one  vessel  into  another  in  the  air  ?  under  water  ?  Could 
hydrogen  be  collected  by  downward  displacement  in  the 
air  ?  by  upward  displacement  ?  Why  not  collect  it  thus 
instead  of  over  water  ?  Why  can  you  not  collect  ammo- 
nia gas  or  hydrochloric  acid  gas  over  water  ? 

(c)  Arrange  an  apparatus  for  burning  hydrogen  in  a 
jet.  Does  the  flame  afford  much  light  ?  much  heat  ? 
What  is  formed  in  the  combustion  ?  Hold  a  dry  beaker 
over  the  hydrogen  jet.  What  is  deposited  on  the  inner 
surface  of  the  beaker  ?  What  remains  in  solution  in  the 
flask  after  making  the  hydrogen?  Write  the  reaction. 
Can  other  metals  and  other  acids  be  used  in  making  hy- 
drogen? Try  several  in  test-tubes.  By  dissolving  a 
metal  in  an  add  a  salt  is  formed. 

OXYGEN. 

We  have  seen  that  mercuric  oxide  when  heated  de- 
composes into  mercury  and  oxygen :  HgO  =  Hg  -f  O. 
Oxygen  can  be  made  also  by  heating  some  other  oxides 
rich  in  oxygen.  Black  oxide  of  manganese  (MnO2) 
when  heated  decomposes  as  follows :  3MnO.2  =  Mn3O4  -f 
O2.  Oxygen  is  generally  made  by  heating  potassium 
chlorate  (KC1O3).  By  mixing  with  the  KC1O3  some 
MnO2  or  Fe2O3  the  oxygen  conies  oif  at  a  lower  tempera- 
ture and  more  regularly.  The  MnO2  should  be  heated 
to  redness  to  insure  the  destruction  of  any  impurity  of 
organic  matter  before  using  it  to  mix  with  KC1O3. 


GENERAL  CHEMISTRY.  25 

31.  Mix  about  equal  weights  of  KC1O3  and  MnO2.     Fill 
a  test-tube  about  one-third  full  of  the  mixture,  heat  and 
collect  the  gas  over  water,  and  try  the  following  exper- 
iments : 

(a)  Plunge  the  glowing  end  of  a  splinter  into  a  bottle 
of  oxygen. 

(5)  Fasten  a  piece  of  charcoal  to  a  large  iron  wire, 
heat  it  to  redness,  and  lower  into  a  bottle  of  oxygen. 

(c)  Ignite  some  sulphur  placed  in  a  hollowed  crayon, 
and  lower  it  into  a  bottle  of  oxygen. 

(d)  Same  as  (c),  substituting  a  piece  of  dry  phosphorus 
for  sulphur. 

(e)  Same  as  (c),  substituting  a  piece  of  sodium  for  sul- 
phur. 

(/)  Attach  a  net  of  fine  iron  wire  to  the  end  of  a  large 
iron  wire,  heat  to  redness,  and  plunge  quickly  into  a  bot- 
tle of  oxygen. 

What  are  formed  in  the  above  combustions  ?  Pour  a 
little  water  into  the  bottles  in  which  the  combustions 
have  been  made,  and  test  the  reaction  with  litmus  paper. 
After  the  test-tube  in  which  the  mixture  was  heated  has 
cooled,  pour  some  water  into  it,  warm,  and  filter.  Add  a 
few  drops  of  silver  nitrate  solution  to  a  portion  of  the 
filtrate.  Add  a  few  drops  of  silver  nitrate  solution  to  a 
solution  of  KC1O3.  The  change  which  the  KC1O,  has  un- 
dergone is  expressed  by  the  equation  KC1O3  =  KC1  +  3O. 

The  reaction  with  nitrate  of  silver  is  as  follows :  KC1  -f 
AgNO3  =  KNOS  -f  AgCl.  AgCl  is  insoluble  in  water. 
Is  it  soluble  in  HNO3  ?  Will  other  chlorides  in  solution 
give  a  precipitate  with  AgNO3  ? 

NITROGEN. 

32.  (a)  Float  a  small  porcelain  crucible  on  the  surface 
of  some  distilled  water  in  a  large  dish.     Dry  by  means 

3    « 


26  EXPERIMENTS   IN 


of  filter  paper  a  piece  of  phosphorus  half  the  size  of  a 
pea,  and  place  it  in  the  crucible.  Set  fire  to  the  phospho- 
rus and  lower  over  it  carefully  a  large-mouthed  bottle. 
When  the  P2O5  fumes  produced  by  the  combustion  have 
been  absorbed  by  the  water,  mark  with  a  label  the  level 
of  the  water  in  the  bottle.  The  gas  remaining  in  the 
bottle  is  nitrogen.  Close  the  mouth  of  the  bottle,  and 
take  it  out  of  the  water. 

(6)  Pour  some  clear  lime-water  into  the  bottle,  and 
shake  it  up  with  the  nitrogen.  Does  nitrogen  render 
lime-water  milky?  Is  nitrogen  combustible?  Does  it 
support  combustion  ?  Measure  the  capacity  of  the  bot- 
tle when  filled  up  to  the  label.  Supposing  the  atmosphere 
to  consist  of  oxygen  and  nitrogen  gases,  what  percentage 
by  volume  is  oxygen  ?  nitrogen  ? 

(c)  Nitrogen  may  be  made  by  heating  a  cone,  solution 
of  ammonium  nitrate  (NH4lSrO2  =  N2  -f  2H2O) ;  or  by 
heating  a  mixture  of  potassium  bichromate  and  ammo- 
nium chloride  (K2O2OT  +  2NH4C1  =  2KC1  +  Cr203  + 
4H20  +  Nf)- 

NITROUS  OXIDE. 

33.  Heat  10  g.  of  ammonium  nitrate  in  a  large  test- 
tube  until  the  salt  melts ;  then  heat  slowly,  but  contin- 
uously. Collect  the  gas  over  warm  water.  Why  not 
over  cold  water  ?  If  any  white  fumes  appear  in  making 
the  gas,  account  for  them.  The  gas  thus  prepared  is  ni- 
trogen monoxide,  or  nitrous  oxide  (N2O).  Write  the  re- 
action. 

(a)  Insert  a  glowing  splinter  into  a  bottle  of  the  gas. 

(6)  Burn  some  sulphur  in  a  bottle  of  the  gas.  Does 
the  sulphur  burn  exactly  as  sulphur  in  oxygen  ? 

(c)  Invert  a  bottle  of  the  gas  in  a  dish  of  cold  water. 
If  the  gas  does  not  all  dissolve,  test  the  residual  gas  with 
a  glowing  splinter. 


GENERAL  CHEMISTRY.  27 

(d)  Incline  a  bottle  of  the  gas  standing  over  water 
until  some  bubbles  of  air  enter.    Is  any  change  apparent  ? 

Save  a  bottle  of  nitrous  oxide  for  34  (e). 

NITRIC  OXIDE, 

34.  Mix  5  cc.  of  water  with  5  cc.  of  cone.   HXO3. 
Pour  this  diluted  acid  over  5  g.  of  copper  turnings  in  a 
small  flask.     The  reaction  is  3Cu  +  8HXO3  =  3Cu(XO3)2 
4-  4H2O  +  2XO  (nitric  oxide,  or  nitrogen  dioxide).     The 
formula   is   also   written    X2O2.     Collect   the   gas   over 
water. 

(a)  When  several  bottles  of  the  gas  have  been  col- 
lected let  the  remainder  of  the  gas  pass  into  a  solution  of 
ferrous  sulphate  (FeSO4).  The  nitric  oxide  is  absorbed 
by  the  solution  to  form  FeSO4XO. 

(6)  Incline  a  bottle  of  nitric  oxide  standing  over  wa- 
ter until  a  few  bubbles  of  air  enter.  Nitrogen  tetroxide 
is  formed,  written  XO2  or  X2O4.  Test  with  litmus  paper 
the  water  in  the  bottle.  2NO2  +  H2O  =  HXO3  + 
HX02.  The  nitrous  acid  (HXO2)  is  at  once  decomposed 
at  ordinary  temperatures  into  nitric  oxide  and  nitric 
acid.  3XO2  +  H2O  =  2HXO3  +  XO. 

Account  for  the  red  fumes  which  were  at  first  seen  in 
the  test-tube  in  preparing  nitric  oxide.  What  became 
of  them  ? 

(c)  Mix  a  small  portion  of  the  dark  solution  (a)  with 
cone.  H2SO4,  keeping  the  mixture  cool. 

(jd)  Pour  the  remainder  of  the  dark  solution  into  a 
test-tube  and  heat.  Is  any  gas  given  off? 

(e)  How  can  oxygen,  nitrous  oxide,  and  nitric  oxide 
be  distinguished  from  one  another. 

NITRIC  ACID. 

35.  (a)  Put  5  g.  of  sodium  nitrate  into  a  large  test- 


28  EXPERIMENTS   IN 


tube;  cover  it  with  cone.  H2SO4;  heat,  collecting  the 
distillate  in  a  test-tube  cooled  in  water.  2NaNO3  + 
H2SO4  =  Na2SO4  +  2IINO3.  Nitric  acid  is  manufact- 
ured from  these  materials. 

(5)  Take  5  cc.  cold  dilute  HNO3  in  a  test-tube.  In- 
cline the  test-tube  and  add  carefully  an  equal  volume  of 
cone,  solution  of  FeSO4,  so  that  the  FeSO4  solution  will 
form  a  layer  above  the  HNO3.  Set  the  tube  aside.  A 
dark  ring  should  form  where  the  liquids  are  in  contact. 

(c)  Add  to  3  cc.  of  a  weak  solution  of  any  nitrate  an 
equal  volume  of  cone.  H2SO4.  Allow  the  mixture  to 
cool.  Now  add  cone,  solution  of  FeSO4  as  in  35  (6). 
What  is  the  test  for  a  nitrate  ?  Why  should  the  mixt- 
ure of  H2SO4  and  the  solution  to  be  tested  for  HNO3  be 
cooled  before  the  addition  of  the  FeSO4  ? 

(d~)  Pour  a  few  drops  of  cone.  HNO3  on  some  pieces  of 
tin  foil  in  a  crucible.  Wash  the  white  powder  obtained 
several  times  with  water,  decanting  the  liquid.  Then 
place  the  white  powder  on  a  filter,  and  wash  until  the 
washings  contain  no  HNO3.  (How  can  the  absence  of 
HNO3  be  proved  ?)  Prove  that  the  white  powder  is  not 
a  nitrate  by  applying  the  test  for  a  nitrate.  All  nitrates 
are  soluble  in  water.  The  tin  is  oxidized  by  the  HNOS. 
4HNO3  =  2H2O  +  2N2O4  +  O2 ;  Sn  +  O2  =  SnO2. 

(e)  Moisten  a  quill  or  other  animal  organic  matter 
with  cone.  HNO3.  What  color  is  imparted  to  it  ?  Will 
ammonia  remove  the  stain  ?  Will  water  wash  it  out  ? 

HYDROCHLORIC  ACID. 

36.  (a)  Put  into  a  300  cc.  flask  20  g.  of  common  salt, 
insert  a  funnel  or  thistle  tube,  which  must  reach  nearly 
to  the  bottom  of  the  flask,  and  connect  the  flask  with 
three  bottles,  A,  B,  and  C,  B  and  C  being  half-full  of 
water.  Mark  with  a  label  the  level  of  the  water  in  B. 


GENERAL   CHEMISTRY.  29 

Pour  40  cc.  of  cone.  H2SO4  in  the  flask,  and  warm  gently. 
Why  do  bubbles  at  first  pass  up  through  the  water  in  B 
and  C,  and  afterward  do  not  appear  in  C  ?  Disconnect 
A,  connecting  the  generating  flask  with  B. 

(6)  Test  the  gas  in  A  by  (1)  holding  near  its  mouth  a 
drop  of  cone,  ammonia  suspended  on  the  end  of  a  glass 
rod,  (2)  inserting  into  the  bottle  a  lighted  splinter,  (3) 
closing  the  mouth  of  the  bottle  and  inverting  it  in  water. 

(c)  Notice  the  level  of  the  water  in  B.     Test  a  portion 
of  the  solution  in  B  by  adding  a  few  drops  of  silver  ni- 
trate.    Is  the  white  precipitate  thus  formed  soluble  in 
HN03? 

(d)  The  reaction  in  making  HC1  is  NaCl  +  H2SO4  = 
NaHSO4  -f  HC1,  or  at  a  higher  temperature  2NaCl  + 
H2SO4  =  ]STa2SO4  +  2HC1.     In  which  of  these  reactions 
is  the  H2SO4  more  economically  used  ?     Both  reactions 
occur  in  making  the  HC1  on  a  large  scale. 

Instead  of  using  a  funnel  tube  in  the  flask  in  the  above 
experiment,  the  bottles  B  and  C  may  each  be  provided 
with  a  third  tube  open  at  both  ends,  passing  through  the 
stopper  and  reaching  below  the  surface  of  the  liquid  in 
the  bottle. 

AQUA  REGIA,  OR  NITRO-HYDROCHLORIC  ACID. 

37.  Mix  2  cc.  of  cone.  HNO,  with  6  cc.  of  HC1.  This 
mixture  is  called  aqua  regia.     It  dissolves  gold.  It  is 
also  an  excellent  solvent  for  many  other  metals.  The 
solution  contains  the  chloride  of  the  metal. 

CARBON  DIOXIDE. 

38.  Pour  a  little  HC1  into  a  test-tube  containing  a 
small  quantity  of  sodium  carbonate.     Apply  a  lighted 
splinter  to  the  mouth  of  the  tube.     Suspend  in  the  tube 
a  drop  of  clear  lime-water  on  the  end  of  a  glass  rod. 


30  EXPERIMENTS. 


The  gas  given  off  is  carbon  dioxide.  Na2CO,  -j-  2HC1  = 
2NaCl  +  H2O  -f  CO2.  Would  the  same  gas  be  evolved 
if  other  carbonates  were  used  ?  other  acids  ?  Try  some 
of  them. 

Put  20  g.  of  marble  (CaCO,)  in  small  pieces  into  a  flask. 
Cover  it  with  HC1.  Collect  the  gas  by  downward  dis- 
placement. 

(a)  Balance  two  beakers  on  a  scale,  and  pour  into  one 
of  them  CO2  from  a  bottle  or  beaker. 

(b~)  Pour  CO2  from  one  beaker  into  another  in  which 
a  wax  taper  is  burning ;  into  a  beaker  containing  lime- 
water.  What  is  formed  when  CO2  is  conducted  into 
lime-water.  Try  the  effect  of  HC1  on  this  precipitate. 

(c)  Breathe  through  a  glass  tube  into  a  beaker  con- 
taining diluted  lime-water.  If  the  lime-water  becomes 
cloudy,  see  if  you  can  make  it  clear  by  continuing  to 
pass  into  it  air  from  the  lungs.  What  does  this  exper- 
iment prove  ? 

(cT)  Conduct  CO2,  passed  through  a  little  water  to  wash 
it,  into  5  cc.  of  blue  litmus  solution  in  a  test-tube  until 
the  color  of  the  solution  is  changed.  Now  warm  the  lit- 
mus solution,  and  observe  the  change  of  color.  What 
does  this  experiment  prove?  State  some  of  the  prop- 
erties of  carbon  dioxide  gas.  What  is  the  test  for  a  car- 
bonate ? 


PART  II. 

EXPERIMENTS  IN  QUALITATIVE  ANALYSIS. 


THE  METALS. 

SEPARATION  OF  THE  METALS  INTO  GROUPS. 

39.  Label  twenty-four  test-tubes  as  follows:  Ag,  Al, 
As',  Ba,  Bi,  Ca,  Cd,  Co,  Cr,  Cu,  Fc,  Hg7,  Hg",  K,  Mg, 
Mn,  Na,  (XH4),  Ni,  Pb,  Sb,  Sn,  Sr,  and  Zn.  Place  the 
tubes  in  a  rack,  and  pour  into  each  of  them  5  cc.  of  a  so- 
lution of  a  salt  of  a  metal  marked  on  the  label. 

(#)  Add  two  or  three  drops  of  HC1  to  the  liquid  in 
each  tube.  If  a  precipitate  forms,  determine  whether  it 
is  soluble  in  a  larger  quantity  of  the  acid.  Group  I.  in- 
cludes those  metals  which  form  chlorides  insoluble  in 
HC1.  Which  are  they  ? 

(6)  Pour  into  each  one  of  the  tubes  which  contains  a 
salt  of  a  metal  not  belonging  to  Group  I.  5  cc.  of  H2S  so- 
lution. Add  a  larger  quantity  of  the  H2S  in  cases  in 
which  there  is  an  indication  that  a  precipitate  is  forming. 
Group  II.  includes  all  those  metals  (exclusive  of 
Group  I.)  which  are  precipitated  as  sulphides  in  the 
presence  of  free  acid.  Which  are  they  ?  Are  the  mem- 
bers of  Group  I.  precipitated  by  H2S  in  presence  of  free 
acid  ? 

(c)  To  the  remaining  solutions  add  NH4OH  to  neutral- 
ize the  free  acid  (What  salt  of  ammonia  is  formed? 
What  does  the  XH4OH  form  with  the  H2S  in  solution?), 


32  EXPERIMENTS  IN 


and  then  a  few  drops  of  yellow  ammonium  sulphide, 
(NH4)2SZ.  If  a  precipitate  forms,  add  more  (NH4)2SX. 
Group  III.  includes  those  metals  which  are  not  pre- 
cipitated by  H2S  in  the  presence  of  dilute  acids,  but  by 
H2S  in  alkaline  solutions,  or  (NH4)2SX  in  neutral  or  al- 
kaline solutions.  Which  metals  belong  to  this  group? 
Are  there  any  members  of  Groups  I.  and  II.  precipitated 
by  H2S  in  alkaline  solutions,  or  (NH4)2SX  in  neutral  or 
alkaline  solutions  ? 

(d)  To  the  remaining  solutions  add  ammonium  car- 
bonate ((]STH4)2CO8)  solution.     If  any  of  the  carbonates 
thus  precipitated  are  soluble  in  NH4C1,  do  not  include 
them  in  Group  IV.     Group  IY.  includes  those  metals 
(exclusive  of  Groups  I.,  II.,  and  III.)  which  form  with 
(NH4)2CO,  carbonates  insoluble  in  water  or  in  NH4Ch  ? 
Name  them.     Must  the  solution  to  which  the  (NH4)2CO3 
is  applied  for  the  separation  of  Group  IY.  have  an  acid 
or  an  alkaline  reaction  ?    Add  (NH4)2CO3  to  an  acidified 
solution  of  a  salt  of  Group  IY.     Does  (NH4)2CO3  precip- 
itate members  of  Groups  I.,  II.,  and  III.  ? 

(e)  Add  sodium  phosphate  (Na2HPO4)  to  the  remain- 
ing solutions.     Group  Y.  is  thus  separated.     What  does 
this  group  include  ?     Does  Na2HPO4  precipitate  members 
of  Groups  I.,  II.,  III.,  and  IY.  ? 

(/)  Group  VI.  embraces  the  metals  still  remaining  in 
solution.  Which  are  they  ? 

Make  out  a  table  for  the  separation  of  a  mixture  of  so- 
lutions of  salts  of  the  above  metals  into  groups. 

The  student  should  now  be  required  to  determine 
what  groups  are  represented  in  a  mixture  given  to  him. 
A  group  must  be  completely  removed  before  the  test  for 
the  following  group  is  applied  (?).  How  can  it  be  deter- 
mined that  the  group  is  completely  removed?  Boiling 
generally  facilitates  the  formation  and  separation  of  a 


QUALITATIVE  ANALYSIS. 


33 


precipitate ;  and  the  filtering  also  proceeds  more  rapidly 
when  the  liquid  is  hot. 

GROUP  I. 
THE  HYDROCHLORIC  ACID  GROUP. 

40.  (a)  Add  HC1  to  solutions  of  lead  nitrate  (Pb(NO3)2), 
silver  nitrate  (AgNO3),  and  mercurous  nitrate  (Hg/NO3). 
Test  the  solubility  of  the  precipitates  in  boiling  water. 
Test  their  solubility  in  ammonia.  Add  HC1  to  the  solu- 
tion in  ammonia  until  the  reaction  is  acid.  Is  all  of  the 
element  reprecipitated  ? 

(6)  Test  solutions  of  Pb(NO3)2,  AgNO3,  Hg/NOs,  and 
mercuric  chloride  (Hg"Cl2)  with  NaOH,  H2SO4,  potas- 
sium iodide  (KI),  potassium  chromate  (K2CrO4),  and 
stannous  chloride  (SnCl2),  and  record  the  results  as  indi- 
cated in  the  following  table  : 


Reagent 
Added  to. 

Pb(N03)2. 

AgNos. 

Hg'No,. 

Hg"Cl2. 

Soluble 
in  (?). 

NaOH 

NaOH 

H2SO* 

NaOH 

KI 

KI 

SnCl2 

K2Cr04 
HC1 

(c)  Heat  in  a  dry  ignition  tube  a  small  quantity  of  a 
mercury  compound ;  a  small  quantity  of  the  same  com- 
pound mixed  with  Na2CO3. 

(d)  Heat  on  charcoal  in  the  reducing  flame  of  the 
blow-pipe  AgCl  (obtained  as  in  a),  PbCl2,  or  PbSO4.     Are 
globules  of  Pb  and  Ag  malleable?     Can  you  cut  them 
with  a  knife  ?     Can  you  mark  on  paper  with  them  ? 

Make  out  a  scheme  (from  reactions  in  a)  for  the  detec- 
tion and  separation  of  Pb,  Ag,  and  Hg7.  What  confirm- 
atory test  could  you  apply  for  each  ? 


34  EXPERIMENTS   IN 


How  can  a  mercurous  be  distinguished  from  a  mer- 
curic salt  ?  an  iodide  from  a  chromate  ? 

The  student  should  now  be  given  exercises  in  the  de- 
tection and  separation  of  the. elements  in  Group  I. 

GROUP  II. 
HYDROGEN  SULPHIDE  GROUP. 

41.  This  group  is  divided  into  Sub-group  A,  metals 
whose  sulphides  are  soluble  in  (NH4)2SX ;  and  Sub-group 
B,  metals  whose  sulphides  are  insoluble  (or  nearly  so)  in 
(NH4)2SX.  Determine  by  experiment  which  metals  be- 
long to  A.  The  precipitated  sulphides  of  Group  II.  have 
generally  the  following  composition :  Sb2S3,  As2S3,  Bi2S3, 
CdS,  CuS,  PbS,  HgS,  SnS  (brown),  and  SnS2  (yellow). 

SUB-GROUP  A. 

Sulphides  Soluble  in  Ammonium  Sulphide. 

(a)  Compare  the  solubility  of  Sn,  Sb,  As,  SnO2,  Sb2O3, 
As2O3,  SnS,  Sb2S3,  and  As2S3  in  boiling  cone.  HC1 ;  in  boil- 
ing cone.  HNO3.  Record  the  results  in  tabular  form. 
(The  sulphides  for  this  experiment  may  be  gotten  by 
precipitating  stannous  chloride  (SnCl2),  antimonious 
chloride  (SbCl3),  and  arsenious  chloride  (AsCl3)  with 
H2S,  and  washing  the  precipitate.)  The  best  solvent  for 
the  metals  is  nitro-hydrochloric  acid  (41  a),  or  cone.  HC1 
to  which  a  crystal  of  KC1O3  is  added,  forming  stannic 
chloride  (SnCl4),  SbCls,  and  arsenic  acid  (H3AsO4). 

(6)  Compare  the  solubility  of  SnS,  Sb2S3,  and  As2S3  in 
boiling  cone.  (NH4)2CO3. 

(c)  Put  a  piece  of  platinum  foil  in  a  porcelain  dish  and 
place  on  it  a  piece  of  zinc.  Cover  the  Zn  and  Pt  with  a 
solution  of  SnCl2.  If  the  Zn  is  not  acted  upon,  acidify 
with  HC1  until  H  is  freely  given  off.  Try  a  similar  ex- 


QUALITATIVE  ANALYSIS.  35 

periment  using  SbCl,  instead  of  SnCl2.  (This  experiment 
must  not  be  tried  with  a  compound  of  arsenic,  nor  in  a  mixt- 
ure of  these  elements  until  any  arsenic  that  may  be  pres- 
ent has  been  separated  from  the  mixture.) 

Touch  the  platinum  foil,  if  it  be  blackened,  with  a 
drop  of  cone.  HXO3  on  the  end  of  a  glass  rod  ;  also  touch 
it  at  another  point  with  a  drop  of  cone.  HC1. 

(d)  Add  Hg"Cl2  to  solutions  of  SnCl2,  SbCl,,  AsCl, 
(40  b). 

(e)  Blow  a  small  bulb  on  the  end  of  an  ignition  tube 
10  cm.  in  length.     Mix  a  small  quantity  of  a  dry  com- 
pound of  arsenic  with  six  times  its  weight  of  a  dry  mixt- 
ture  of  equal  parts  of  Xa2COs  and  potassium  cyanide 
(KCX),  place  in  the  bulb,  and  heat.     The  bulb  should 
not  be  more  than  half-full  of  the  mixture.     If  a  mirror 
of  arsenic  forms  in  the  upper  part  of  the  tube,  break  the 
tube  and  test  the  solubility  of  the  arsenic  in  sodium  hy- 
pochlorite  (NaOCl).     A  mirror  of  the  Sb  is.not  soluble 
in  this  reagent. 

(/)  Mix  a  very  small  quantity  of  a  compound  of  arsenic 
with  Na2CO3  and  KCN,  and  heat  on  charcoal  in  the  re- 
ducing flame  of  the  blow-pipe.  Notice  the  odor  of  the 
fumes.  Heat  in  like  manner  an  Sb  compound,  and  when 
a  globule  is  obtained  throw  the  melted  globule  from  the 
charcoal  on  the  floor.  Let  another  globule  of  Sb  cool, 
and  compare  its  color,  solubility,  and  malleability  with 
the  corresponding  properties  of  a  globule  of  Sn  gotten 
in  the  same  way. 

(g)  Compare  the  action  of  H2S  on  an  arsem'ous  and  an 
arsenic  compound ;  on  a  stan/iows  and  a  stannic  com- 
pound. 

(h)  Make  some  "  Magnesia  Mixture  "  by  precipitating 
MgSO4  solution  with  a  large  excess  of  NH4OH,  then  add- 
ing NH4C1  to  the  precipitate  and  solution  until  the  pre- 


36  EXPERIMENTS   IN 


cipitate  is  dissolved.  Add  "Magnesia  Mixture"  to  so- 
dium arsenite ;  to  sodium  arsenate ;  to  sodium  phosphate. 
Add  AglSTOg  to  a  neutral  solution  of  sodium  arsenite ;  to 
sodium  arsenate ;  to  sodium  phosphate.  Add  ammonium 
molybdate  solution  to  a  few  drops  of  sodium  arsenite ;  to 
sodium  arsenate ;  to  sodium  phosphate ;  and  if,  after  a 
few  minutes,  no  precipitate  occurs,  warm.  How  can  ar- 
senites,  arsenates,  and  phosphates  be  distinguished  from 
each  other  ? 

Can  the  precipitation  of  Sn,  Sb,  and  As  by  H2S  be  pre- 
vented by  the  presence  of  strong  HC1  ?  Which  of  these 
sulphides  most  readily  dissolves  in  HC1  ?  (41  6).  Make 
out  a  scheme  for  the  detection  and  separation  of  Sn,  Sb, 
and  As. 

An  exercise  should  be  given  in  Sub-group  A. 

SUB-GROUP  B. 

Sulphides  Insoluble  in  Ammonium  Sulphides. 

(a)  Which  members  of  Sub-group  B  are  precipitated 
from  solutions  by  H2SO4?  by  NH4OH?  Which  sul- 
phides of  Group  II.  B  are  dissolved  by  boiling  with 
HNO3?  by  boiling  with  H2SO4?  by  KCN  solution? 
Make  out  the  results  of  these  experiments  in  tabular 
form.  From  the  above  reactions  make  out  a  scheme  for 
the  separation  of  the  members  of  Sub-group  B. 

(6)  Dissolve  the  sulphide  which  is  not  soluble  in  HKO, 
in  cone.  HC1  to  which  a  crystal  of  KC1O,  is  added,  boil 
off  the  chlorine  and  apply  a  test  already  known  for  the 
metal  in  solution. 

(c)  Add  NH4OH  to  a  solution  of  BiCl3.  The  precip- 
itate is  Bi(OH)2.  Dissolve  this  precipitate  in  as  small  a 
quantity  of  cone.  HC1  as  possible.  The  solution  contains 
bismuth  chloride  (BiCl3).  Dilute  with  water.  Bismuth 
oxychloride  (BiOCl)  is  precipitated.  Dilute  with  water 


QUALITATIVE  ANALYSIS.  37 

solutions  of  BiCl3  and  SbCl3.  Is  either  precipitate  solu- 
ble in  tartaric  acid  ?  The  atmosphere  sometimes  precip- 
itates stannous  oxychloride  (SnOCl)  from  solutions  of 
SnCl2  on  standing.  How  could  this  be  distinguished 
from  the  oxychlorides  of  Bi  and  Sb  ? 

(d)  Heat  Bi2Ss  on  charcoal  before  the  blow-pipe.     Test 
in  like  manner  CdS. 

(e)  Sulphur  formed  by  the  decomposition  of  H2S  or 
(HN4)2SX  may  be  mistaken  for  a  metallic  sulphide.     It 
may  be  recognized  by  the  tests  (28  a).     Add  an  acid  to 
H2S.    Add  an  acid  to  (NH4)&.    Add  H2S  to  K2CrO4. 
Add  H2S  to  Fe2Cl6.     CuS  is  slightly  soluble  in  warm 
(NH4)2SX,  but  not  in  sodium   hydrosulphide  (NaSH) ; 
but  HgS  is  somewhat  soluble  in  the  latter.     Hence  in 
the  absence  of  Hg,  NaSH  is  sometimes  used  instead  of 
(NH4)2SX  for  the  separation  of  the  sulphides  of  Sub-group 
B  from  those  of  Sub-group  A.     Sodium  hydrosulphide  is 
made  by  saturating  a  solution  of  NaOH  with  H2S. 

An  exercise  should  be  given  in  Sub-group  A ;  also  one 
in  a  mixture  of  Sub-groups  A  and  B,  or  in  a  mixture  of 
Groups  I.  and  II. 

GROUP  III. 
AMMONIUM  SULPHIDE  GROUP. 

42.  (a)  Heat  Co,  Mn,  and  Ni  compounds  in  the  borax 
bead.  Heat  a  Mn  compound  with  a  mixture  of  Na2CO$ 
and  Naj^O,  on  platinum  foil.  Heat  a  Zn  compound 
moistened  with  cobalt  nitrate  on  platinum  foil.  Heat 
an  Al  compound  strongly  on  charcoal,  moisten  with  a 
solution  of  cobalt  nitrate  and  heat  again.  Try  a  like  ex- 
periment with  calcium  phosphate  (CasPO4). 

(6)  Place  in  test-tubes  5  cc.  each  of  the  nitrates,  chlo- 
rides, or  sulphates  of  Al,  Co,  Cr,  Fe',  Fe",  Mn,  Ni,  Zn. 
Add  NH4OH.  The  hydroxides  are  formed,  A12(OH)6) 


38  EXPERIMENTS   IN 


Co(OH)2,  Cr2(OH)6,  Fe(OH)2,  Fe2(OH)6,  Mn(OH)2, 
Ni(OH)2,  Zn(OH)2.  Which  are  dissolved  in  excess  of 
NH4OH  ?  Which  are  soluble  in  NH4C1,  and  consequent- 
ly not  precipitated  in  the  presence  of  a  sufficient  quan- 
tity of  NH4C1?  Is  Fe(OH)2  at  all  soluble  in  NH4OH? 

(c)  Test  FeSO4  and  Fe2Cl6  separately  with  NH4OH, 
with  potassium  ferrocyanide  (K4FeCy6),  and  with  potas- 
sium sulphocyanate  (KCNS).     Boil  FeSO4  with  a  few 
drops  of  strong  HNOS  until  the  solution  changes  color ; 
then  test  portions  of  it  with  NH.OH,  K4FeCy6,  KCNS. 

(d)  Add   solution   of  lead   acetate   to    a   solution   of 
K2CrO4.     Add  solution  of  AgNO3  to  a  solution  of  K2CrO4. 
Mix  Cr2(OH)6  (see  its  formation  above)  with   NaNO, 
and  Na2CO3;  treat  the  fused  mass  with  water,  acidify 
with  acetic  acid,  and  test  the  solution  with  lead  acetate ; 
with  AgNO,.     How  can  a  compound  insoluble  in  water 
be  tested  for  Or? 

(e)*To  solutions  of  nitrate,  chloride,  or  sulphate  of  Al, 
Cr,  Fe'  add  slowly  NaOH.  The  hydroxides  are  precip- 
itated. Which  of  these  precipitates  dissolve  in  excess  of 
NaOH  in  the  cold?  Do  any  dissolve  on  boiling?  Are 
any  that  dissolve  in  the  cold  reprecipitated  by  boiling  ? 
by  adding  NH4C1  ?  How  can  Al,  Cr,  and  iron  in  a  mixt- 
ure be  separated,  and  the  presence  of  each  be  shown, 
the  iron  being  in  a  ferric  condition  ?  the  iron  being  in  a 
ferrous  condition  ?  Try  the  latter. 

(/)  To  nitrates,  chlorides,  or  sulphates  of  Co,  Mn,  Ni, 
and  Zn  add  (NH4)2SX.  The  precipitates  are  CoS,  MnS, 
NiS,  and  ZnS.  Note  their  color.  Which  are  soluble  in 
HC1  ?  acetic  acid  ?  nitro-hydrochloric  acid  ? 

(g)  Add  NaOH  to  solutions  of  Co,  Mn,  Ni,  and  Zn. 
The  hydroxides  Co(OH)2,  Mn(OH)2,  Ni(OH)2,  Zn(OH), 
are  formed.  Are  any  of  these  precipitates  soluble  in 


QUALITATIVE  ANALYSIS.  39 

(A)  To  5  cc.  cobaltous  chloride  (CoCl2)  add  solution 
of  KCX  gradually.  The  precipitate  is  cobaltous  cyanide 
(Co(CX)2).  On  adding  an  excess  of  KCX  the  precipi- 
tate dissolves,  and  the  solution  contains  the  double  salt, 
potassium  cobalto-cyanide  (K4Co(CX)6),  analogous  to  po- 
tassium ferro-cyanide  (K,Fe(CX)6) :  Co(CN)2  +  4KCN 
=  K4Co(CX)6.  Add  one  or  two  drops  of  HC1  to  the  solu- 
tion, heat  it  nearly  to  the  boiling  point,  and  allow  it  to 
stand  for  five  minutes.  Then  to  a  portion  of  the  solution 
add  excess  of  HC1 ;  to  the  other  portion  add  NaOH  in 
large  excess,  warm  gently,  and  add  bromine  water. 

(i)  The  same  as  (A)  using  nickelous  chloride  (NiCl2) 
instead  of  CoCl2. 

In  (A)  if  instead  of  one  or  two  drops  of  HC1  the  solu- 
tion be  neutralized  or  acidified  without  heating,  the  potas- 
sium cobalto-cyanide  is  decomposed  and  Co(CN)2  is  pre- 
cipitated. But  when  only  a  few  drops  of  the  acid  are 
added  (the  KCN  remaining  in  excess)  and  the  solution 
is  heated  the  free  hydrocyanic  acid  (HCX)  converts 
the  Co(CN)2  into  cobaltic  cyanide  (Co2(CX)6),  which 
unites  with  KCN  to  form  potassium  cobalti-cyanide 
(K8Co(CX)8),  analogous  to  potassium  ferri  -  cyanide 
(K§Fe(CN)6).  KCX  +  HC1  =  KC1  +  HCX ;  2HCN 
+  2Co(CN)2  =  H2  +  Co^CN).;  Co2(CX)6  +  6KCN  = 
K6Co,(CN)12,  or  2K,Co(CN)0.  K,Co(CN).  is  not  decom- 
posed by  dilute  acids  and  is  not  precipitated  by  wanning 
with  excess  of  XaOH  and  adding  bromine  water. 

Potassium  nickelous-cyanide  (K4Ni(CN)6)  in  presence 
of  KCN  is  not  converted  by  the  action  of  HC1  into  a 
compound  corresponding  to  potassium  cobalto-cyanide, 
but  is  decomposed,  and  Ni(CN)2  precipitated  (?).  When 
XaOH  in  large  excess  is  added  to  a  solution  of  Ki]S'i(C]N")6 
and  the  mixture  is  then  warmed  gently  and  bromine  wa- 
ter added,  nickelic  hydroxide  (Ni(OH)3)  is  precipitated  (?). 


40  EXPERIMENTS   IN 


6  4-  13Br  +  3KOH  =  7KBr  +  Ni(OH)3  + 
6BrCN.  The  precipitate  should  be  tested  for  Ni  by 
means  of  the  sodium  metaphosphate  bead. 

CoS  and  NiS  are  converted  into  CoCl2  and  MC12  by 
dissolving  in  nitrohydrochloric  acid.  Before  adding 
KCN  evaporate  the  above  solution  almost  to  dryness 
and  dissolve  in  water,  adding  a  few  drops  of  HC1  if 
necessary  for  solution. 

Take  a  mixture  of  CoCl2  and  NiCl2,  precipitate  with 
(NH4)2SX,  and  see  if  you  can  prove  the  presence  of  Co 
and  Ni  by  the  above  process. 

Make  out  a  scheme  for  the  separation  and  detection  of 
Al,  Co,  Cr,  Fe',  Fe",  Mn,  Ni,  and  Zn  in  mixtures. 

An  exercise  should  be  given  in  Group  III. 

PRECAUTIONS  IN  THE  SEPARATION  OF  GROUPS  I., 
II.,  AND  III. 

43.  Perform  the  following  experiments,  and  then  state 
what  precautions  are  necessary  in  the  separation  of  the 
members  of  Group  III.  from  the  filtrate  of  Group  II. 

(a)  Boil  1  cc.  Fe2Cl6  with  H2S,  and  then  determine 
whether  the  iron  is  in  a  ferrous  or  a  ferric  state  (42  c). 

(b*)  Boil  a  solution  of  H2S ;  test  the  gas  coming  off 
with  a  piece  of  paper  moistened  with  lead  acetate. 
When  paper  moistened  with  lead  acetate  is  not  black- 
ened by  holding  it  to  the  mouth  of  the  test-tube  test  the 
solution  with  lead  acetate  (39  c). 

(c)  Add  NH4OH  to  a  solution  of  MgSO4.  Add  NH4C1 
to  a  solution  of  MgSO4  and  then  add  NH4OH  (39  c). 

GROUP  IV. 
AMMONIUM  CARBONATE  GROUP. 

44.  (a)  What  colors  do  salts  of  Ba,  Ca,  Sr,  and  Cu  give 
when  moistened  with  HC1  and  held  in  the  flame  on  the 
end  of  a  platinum  wire  ? 


QUALITATIVE   ANALYSIS. 


41 


(ft)  To  solutions  of  BaCl2,  CaCl2,  and  SrCl2  add  the  fol- 
lowing reagents  (test  the  delicacy  of  the  reaction  when 
a  precipitate  is  given  by  diluting  to  a  large  volume) : 


Reagent.    |      BaCl2. 

CaCl2. 

SrCl2. 

CaS04. 

H2SO*. 

Completely 
precipitated 
(?)• 

H2SO4  +  al- 
cohol. 

Hydro  -  fluo- 
silicic    acid 
(H2SiF6). 

' 

H2SiF6-f-  al- 
cohol. 

NH4OH    + 
NH4C1     + 
Ammonium 
o  x  a  1  a  t  e 

((NH4)2C204) 

Is  the  test 
most  deli- 
cate for  Ba, 
Ca,  orSr? 

K2OO4. 

Soluble  in  ace- 
tie  acid  (?). 

Are  Mg  salts  precipitated  by  the  reagents  used  in  the 
above  table?  Are  Ba,  Ca,  and  Sr  completely  precip- 
itated by  NH4OH,  XH4C14,  and  (NH4)2CO8  from  solu- 
tions of  their  salts? 

How  can  Ba,  Ca,  and  Sr  be  completely  removed  from 
a  solution  containing  Ba,  Ca,  Sr,  and  Mg,  leaving  the  Mg 
in  solution? 

(c)  Test  the  solubility  of  finely  powdered  chlorides 
and  nitrates  of  Ba,  Ca,  and  Sr  in  absolute  alcohol,  cold 
and  warm.  Use  very  small  quantities.  The  chlorides 
4 


42  EXPERIMENTS   IN 


and  nitrates  may  be  made  from  the  precipitated  carbon- 
ates by  dissolving  the  carbonates  in  as  little  HC1  or 
HNO3  as  possible,  evaporating  to  dryness,  taking  up  in 
a  little  water  and  evaporating  again. 

(d)  Are  members  of  Group  I.,  II.,  and  III.  precipitated 
by  im4Cl,  £TH4OH,  and  (NH4)2CO3  ?  Make  out  a  scheme 
for  the  separation  and  detection  of  Ba,  Ca,  and  Sr  in 
mixtures. 

An  exercise  should  be  given  on  Group  IV. 

OROUF  V. 
SODIUM   PHOSPHATE   GROUP. 

45.  (a)  Heat  a  magnesium  salt  on  charcoal  before  the 
blow-pipe ;  moisten  with  a  solution  of  Co(NO8)2,  and  heat 
again. 

(ft)  To  a  solution  of  magnesium  sulphate  (MgSO4)  add 
NH4C1,  NH4OH,  and  sodium  phosphate  (Na2HPO4). 
Try  this  test  with  a  very  dilute  solution  of  MgSO4. 
Eubbing  the  inside  of  the  test-tube  with  a  glass  rod  will 
facilitate  the  separation  of  the  precipitate.  The  precip- 
itate is  Mg(NH4)PO4. 

(c)  Add  barium  hydroxide  Ba(OH)2  to  a  solution 
of  MgCl2.  The  precipitate  is  magnesium  hydroxide 
(Mg(OH)2).  It  is  somewhat  soluble  in  ammonium 
salts  (?).  What  remains  in  solution  ? 

OROUP  VI. 
SOLUBLE   GROUP. 

46.  (a)  Test  salts  of  JSTH4,  K,  and  Na  on  a  platinum 
wire  in  the  flame.     Examine  the  flame  in  each  case  by 
looking  through  a  piece  of  blue  glass.     Mix  the  three 
compounds  and  examine  in  the  flame  without  the  glass ; 
through  the  glass. 


QUALITATIVE  ANALYSIS.  43 

(6)  Are  NH4,  K,  and  Na  salts  volatile  at  a  faint  red 
heat  ?  Heat  them  on  platinum  foil. 

(c)  Eub  salts  of  NH4,  Xa,  and  K  with  caustic  lime 
moistened  with  water.  Notice  the  odor. 

(cT)  Add  one  or  two  drops  of  platinum  chloride 
(PtCl4)  to  neutral  or  acid  solutions  of  NH4,  Na,  and  K 
salts  on  watch-glasses.  Examine  the  precipitates  with 
a  microscope. 

(e)  Add  tartaric  acid  to  neutral  solutions  of  ]STH4,  Na, 
and  K  salts ;  to  alkaline  solutions. 

(/)  Prove  the  presence  of  NH4,  K,  and  Na  in  a  mixt- 
ure of  solutions  of  NH4C1,  KC1,  and  Nad. 

SEPARATION   OF  GROUPS  IV.,  V.,  AND  VI. 

47.  Take  a  mixture  of  BaCl2,  MgCL2,  and  KC1.  Add 
NH4C1,  NH4OH,  and  (NH4)2CO3.  .  Filter.  To  filtrate 
add  a  few  drops  of  H2SO4.  What  is  the  slight  precip- 
itate ?  Filter.  What  remains  in  solution  in  the  filtrate  ? 
Remove  ammonium  salts  from  the  filtrate  by  evaporating 
and  heating  to  faint  redness.  Dissolve  the  residue  in 
water  and  add  excess  of  Ba(OH)2.  What  is  the  precip- 
itate? What  remains  in  solution?  Add  NH4OH  and 
(XH4)2CO8.  What  is  the  precipitate  ?  Filter.  What  is 
in  the  filtrate  ?  Evaporate  the  filtrate  and  heat  to  red- 
ness. For  what  purpose  ?  Test  the  residue  in  the  flame. 
Dissolve  the  residue  in  water,  adding  a  few  drops  of  HC1 
to  convert  any  K2CO3  that  may  have  been  formed  into 
KC1.  Add  PtCl4  to  the  solution.  Examine  the  precip- 
itate with  a  microscope. 

Make  out  a  scheme  for  the  separation  of  Mg,  K,  and 
Na  from  Group  IY.  and  for  proving  their  presence. 

An  exercise  should  be  given  with  a  mixture  of  salts  of 
Groups  IY.,  Y.,  and  YI.  An  exercise  should  be  given 
with  a  mixture  containing  any  or  all  of  the  Groups. 


44  EXPERIMENTS  IN 


SOME  SPECIAL  CASES  UNDER  GROUP  III. 

Ba,  Ca,  Sr  may  be  precipitated  in  Group  III.  when 
present  in  the  form  of  phosphate,  silicate,  fluoride,  borate, 
or  oxalate.  As  an  example  of  this  we  give  calcium  phos- 
phate (Ca3(P04)2). 

48.  Dissolve  5  g.  Ca3(P04)2  in  HC1.     Divide  the  solu- 
tion into  three  portions,  a,  b,  and  c.     To  a  add  NH4OH 
to  alkaline  reaction.     Dissolve  the  precipitate  in  HNO, 
and  test  for  H8PO4  with  ammonium  molybdate.     Dilute 
b  with  ten  volumes  of  water.     To  a  small  portion  of  this 
solution  add  H2SO4,  then  three  volumes  of  alcohol.     This 
suggests  a  method  for  the  detection  of  Ca,  Ba,  and  Sr 
when  precipitated  as  phosphates  in  Group  III.     Another 
method  is  indicated  by  the  following  experiment :  Al- 
most neutralize  c  with  NaOH ;  add  an  equal  volume  of 
sodium  acetate,  and  then  a  few  drops  of  acetic  acid. 
ISTow  add  Fe2Cl6  drop  by  drop  as  long  as  a  precipitate 
forms  and  until  the*  liquid  after  shaking  remains  reddish. 
Heat  to  boiling  and  filter  hot.     The  precipitate  should 
contain  the  H3PO4  as  phosphate  of  iron  ;  the  filtrate,  the 
Ca  as  CaCl2.     Test  the  precipitate  for  H3PO4  and  the  fil- 
trate for  Ca. 

THE  ACIDS. 

INORCANIC  ACIDS. 
BARIUM    CHLORIDE   GROUP. 

49.  Add  BaCl2  to  the  following  acids  made  neutral 
with  NH4OH,  or  to  neutral  solutions  of  soluble  salts  con- 
taining these  acids,  and  determine  which  of  the  precip- 
itates are  soluble  in  HC1 :  acetic,  arsenious,  arsenic,  ben- 
zoic,    boric,   carbonic,   chloric,   chromic,    citric,    formic, 
hyposulphurous,    hyposulphuric    (hydrogen    sulphide), 
hydrobromic,  hydrochloric,  hydriodic,  hydro ferrocyanic, 


QUALITATIVE  ANALYSIS.  45 

hydroferricyanic,  hydrofluoric,  nitric,  oxalic,  phosphoric, 
sulphurous,  sulphuric,  and  tartaric. 

Of  the  above  acids  acetic,  benzoic,  carbonic,  oxalic,  and 
tartaric  are  organic. 

SILVER   NITRATE  GROUP. 

50.  (a)  Add  a  few  drops  of  AgNO3  to  neutral  solu- 
tions of  salts  of  the  acids  mentioned  in  49.     Xote  the 
color  of  each  precipitate.     Which  of  the  precipitates  are 
soluble  in  HNOS  ?  in  XH4OH  ?  in  excess  of  the  solution 
to  which  the  AgNO3  is  added  ?     Compare  the  results  of 
49  and  50  (a). 

(6)  Put  a  crystal  of  KC1  in  a  test-tube,  add  a  little 
black  oxide  of  manganese  (MnO2),  and  then  a  few  drops 
of  cone.  H2SO4.  Heat  cautiously.  Notice  the  color  of 
the  gas  that  is  given  off.  Try  what  effect  the  gas  has 
upon  the  color  of  some  starch  paste  placed  on  the  end  of 
a  glass  rod,  or  dissolved  in  water.  Test  KBr  and  KI  in 
the  same  way  as  KC1. 

(c)  Test  potassium  ferrocyanide  with  FeSO4  solution ; 
with  a  solution  of  Fe.2Cl6.  Test  potassium  ferricyanide 
with  FeSO4  solution ;  with  a  solution  of  Fe2Cl6.  Add  a 
few  drops  of  Fe2Cl6  to  a  solution  of  FeSO4;  then  add 
some  XaOH,  and  warm.  Is  the  precipitate  soluble  in 
HC1?  To  a  solution  of  potassium  cyanide  (KCN)  add 
FeSO4,  a  few  drops  of  Fe2Cl6,  ISTaOH,  warm ;  then  add 
HC1. 

SILICIC  ACID   (H4SiO4). 

51.  Acidulate  a  solution  of  sodium  silicate  with  HC1, 
and  evaporate  to  dryness  on  the  water  bath.     Treat  the 
residue  with  water  and  HC1,  and  again  evaporate  to  dry- 
ness.    Silica  (SiO2)  insoluble  in  HC1  should  remain.    Heat 
a  particle  of  SiO2  in  the  sodium  metaphosphate  bead. 

Insoluble  silicates  can  be  converted  into  alkaline  sil- 


46  EXPERIMENTS   IN 


icates  by  fusing  with  four  parts  of  mixed  carbonates  of 
potassium  and  sodium. 

HYDROFLUORIC  ACID   (HF). 

52.  Coat  a  watch-glass  with  paraffin.     Trace  a  figure 
in  the  paraffin,  cutting  the  lines  down  to  the  glass.     Mix 
1  g.  CaF2  in  a  lead  crucible  with  cone.  H2SO4  to  form  a 
thin  paste.     Cover  the  crucible  with  the  coated  watch- 
glass,  and  let  it  stand  for  half  an  hour  in  a  warm  place. 
Remove  the  wax  from  the  watch-glass.     The  figure  is 
etched  on  the  glass  (?). 

BORIC  ACID   (H,BO8). 

53.  (a.)  With  a  solution  of  a  borate  form  precipitates 
by  means   of  BaCl2,  CaCl2,  AgNO3,  and  lead  acetate. 
What  are  these  precipitates  soluble  in  ? 

(6)  Mix  some  borax  in  a  porcelain  dish  with  cone. 
H2SO4 ;  cover  it  with  alcohol  and  ignite.  Stir  the  mixt- 
ure while  it  is  burning  and  notice  the  color  imparted  to 
the  flame.  Try  a  similar  experiment,  omitting  the  H2SO4. 
Vary  the  experiment  by  using  boric  acid  instead  of  borax, 
and  omitting  the  H2SO4.  Moisten  borax  with  cone. 
H2SO4  and  heat  in  the  loop  of  a  platinum  wire  in  the 
oxidizing  flame.  Let  the  bead  cool ;  moisten  it  with 
glycerin,  and  heat  again.  What  other  substances  color 
the  flame  green  ? 

(c)  Make  a  solution  of  a  borate  slightly  acid  with  HC1, 
dip  into  it  a  slip  of  turmeric  paper,  and  dry  the  paper  at 
a  gentle  heat. 

SULPHUROUS  ACID   (H2SO3);    HYPOSULPHUROUS 
(THIOSULPHURIC)  ACID   (H2S2O3). 

54.  To  solutions  of  sodium  sulphite  (2sTa2SO3)  and  so- 
dium hyposulphite  (Na2S2O3  add  (1)  HC1;   (2)  AgNO, 
gradually;  (3)  Fe2Cl6;  (4)  CaCl2. 


QUALITATIVE  ANALYSIS.  47 


55.  (a)  Heat  a  crystal  of  KC1O3  before  the  blow-pipe 
on  charcoal. 

(b)  Mix  a  very  small  quantity  of  KC1O,  with  KCN, 
and  heat  on  platinum  foil. 

(c)  Let  fall  a  few  drops  of  cone.  H-jSC^  on  a  crystal  of 
KC1O,  in  a  test-tube.     Warm  gently.     (Caution.) 

ORQANIC  ACIDS. 
TEST  OF  AN    ORGANIC  COMPOUND. 

56.  (a)  The  blackening  of  a  substance  on  heating  in 
an  ignition  tube  usually  indicates  the  presence  of  an  or- 
ganic acid.     Heat  in  this  way  potassium  tartrate,  potas- 
sium citrate,  sodium  acetate. 

(b)  Blackening  without  charring  may  be  due  to  the 
formation  of  a  black  oxide,  and  may  take  place  when  no 
organic  acid  is  present.  Heat  Cu(NO3)2  in  an  ignition 
tube  or  on  platinum  foil. 

The  presence  of  certain  organic  acids  prevent  or  re- 
tard the  precipitation  of  some  of  the  bases  by  the  group 
reagents ;  hence  it  is  best  in  making  an  analysis  to  remove 
or  destroy  organic  acids  by  combustion  as  above,  or  by 
other  methods  given  in  works  on  qualitative  analysis, 
before  applying  the  group  reagents. 

The  organic  acids  are  very  numerous.  The  exper- 
iments here  will  be  limited  to  oxalic,  tartaric,  citric, 
malic,  succinic,  benzoic,  formic,  and  acetic  acids. 

CONG.  H2S04  ON  ORGANIC  ACIDS. 

57.  Pour  cone.  H2SO4  on  small  quantities  of  salts  of 
the  above  acids  in  test-tubes.     Observe  what  takes  place. 
After  a  few  minutes,  heat.     Oxalic  and  formic  acids  are 
decomposed  as  follows :  C2H2O4  =  H2O  -f  CO2  +  CO ; 


48  J  '   *'  |  o "  Vt J  ;    ,*;.  EXPERIMENTS. 


CH2O2  =  H2O  -f  CO.     Carbonic  oxide  (CO)  burns  with 
a  blue  flame.     Test  for  it  at  the  mouth  of  the  test-tube. 

OXALIC,  TARTARIC,   CITRIC,  AND   MALIC  ACIDS, 

58.  To  neutral  solutions  of  salts  of  these  acids  add  : 

(a)  CaCL,.  The  precipitates  are  soluble  in  acetic 
acid  (?).  The  precipitates  dissolve  in  cold  KOH  (?),  and 
fall  again  on  boiling  the  solution  (?). 

(6)  Lime-water  in  considerable  quantity  ;  if  no  pre- 
cipitate appears,  boil. 

(c)  Lead   acetate.     The    precipitates    are    soluble    in 


SUCCINIC,  BENZOIC,  FORMIC,  AND  ACETIC  ACIDS. 

59.  To  solutions  of  salts  of  these  acids  add  : 

(a)  Fe2Cl6  ;  boil  ;  treat  the  precipitates  with  1STH4OH  ; 
then  add  cone.  HC1  in  excess. 

(6)  A  mixture  of  alcohol,  NH4OH,  and  CaCl2. 

(e)  HgNO3;  boil. 

(d)  Warm  a  crystal  of  sodium  acetate  with  cone. 
H2SO4  and  an  equal  volume  of  alcohol.  Acetic  ether  is 
formed.  Notice  the  odor  of  it. 

ORGANIC  COMPOUNDS  INSOLUBLE  IN  WATER. 

To  detect  an  organic  acid  in  a  compound  insoluble 
in  water  the  compound  is  first  decomposed  by  boiling 
with  a  solution  of  Na2CO3.  The  organic  acid  goes  into 
solution  in  the  form  of  a  sodium  salt. 

60.  Boil  calcium  oxalate  (obtained  by  adding  ammo- 
nium oxalate  to  CaCl2)  with  Na2CO3  solution.     Filter. 
Neutralize  the  filtrate  with  acetic  acid,  and  test  for  ox- 
alic acid  with  CaCl.2.     Why  is  acetic  acid  used  instead  of 
HC1?     How  is  calcium  oxalate  distinguished  from  cal- 
cium tartrate  ?  from  calcium  phosphate  ? 


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